Chapter 1. ORGANIC CHEMISTRY STUDYOFCARBON CONTAINING COMPOUNDS Compounds from Nature Synthetic compounds: invented by organic chemists and prepared in
ORGANIC CHEMISTRY STUDYOFCARBON CONTAINING COMPOUNDS Compounds
from Nature Synthetic compounds: invented by organic chemists and
prepared in their laboratories Friedrich Woehlers urea synthesis
Ammonium isocyanate + heat ------> urea NH 4 CNONH 2 CONH 2 I
have been able to make urea without aid of kidney of man or dog.
1828
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Some organic chemicals DNA Essential oils Medicines Active
Pharmaceutical Ingredients Excipients Materials Fuels Pigments
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WITH ITSELF No limit AND
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Electronic Structure of Atoms Structure of atoms a small dense
nucleus, diameter 10 -14 - 10 -15 m, which contains positively
charged protons, neutrons and most of the mass of the atom
extranuclear space, diameter 10 -10 m, which contains negatively
charged electrons
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Notice: one s orbital in each principal shell three p orbitals
in the second shell (and in higher ones) five d orbitals in the
third shell (and in higher ones)
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Rules for Electron Configurations Capacities of shells (n) and
subshells (l)
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Electronic Structure of Atoms Electrons are confined to regions
of space called principle energy levels (shells) each shell can
hold 2n 2 electrons (n = 1, 2, 3, 4......)
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Electronic Structure of Atoms Shells are divided into subshells
called orbitals, which are designated by the letters s, p,
d,........ s (one per shell) p (set of three per shell 2 and
higher) d (set of five per shell 3 and higher).....
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Electronic Structure of Atoms Rule 1: Rule 1: orbitals fill
from lowest energy to highest energy Rule 2: Rule 2: only two
electrons per orbital, spins must be paired Rule 3: Rule 3: for a
set of orbitals with the same energy, add one electron in each
before a second is added in any one
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Periodic Behavior of Elements Flame tests: elements with low
first ionization energies are excited in a flame, and often emit in
the visible region of the spectrum Atoms emit energy when electrons
fall from higher to lower energy states BaSrCa KNaLi
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Atomic Spectrum of Hydrogen
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Electronic Structure of Atoms The pairing of electron
spins
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Lewis Structures Gilbert N. Lewis Valence shell Valence shell:
the outermost electron shell of an atom Valence electrons Valence
electrons: electrons in the valence shell of an atom; these
electrons are used in forming chemical bonds Lewis structure Lewis
structure the symbol of the atom represents the nucleus and all
inner shell electrons dots represent valence electrons For Nitrogen
atom: Valence shell of Nitrogen= 3 Number of valence electrons of
Nitrogen = 5
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Lewis Structures Lewis structures for elements 1-18 of the
Periodic Table For Nitrogen atom: Valence shell of Nitrogen= 3
Number of valence electrons= 5
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Lewis Model of Bonding Atoms bond together so that each atom in
the bond acquires the electron configuration of the noble gas
nearest it in atomic number anion an atom that gains electrons
becomes an anion cation an atom that loses electrons becomes a
cation Ionic bond Ionic bond: a chemical bond resulting from the
electrostatic attraction of an anion and a cation Covalent bond
Covalent bond: a chemical bond resulting from two atoms sharing one
or more pairs of electrons We classify chemical bonds as ionic,
polar covalent, and nonpolar covalent based on the difference in
electronegativity between the atoms
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Electronegativity Electronegativity Electronegativity: a
measure of the force of an atoms attraction for the electrons it
shares in a chemical bond with another atom Pauling scale increases
from left to right within a period increases from bottom to top in
a group
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Electronegativity Electronegativity of atoms (Pauling
scale)
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Electronegativity Electronegativity and chemical bonding
Example: NaCl Na = 0.8, Cl = 3.0 Difference is 2.2, so this is an
ionic bond!
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Coulombs Law The energy of interaction between a pair of ions
is proportional to the product of their charges, divided by the
distance between their centers
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What forces that hold atom together within molecules?
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Covalent Bonding Forces Electron electron repulsive forces
repulsive forces Proton proton repulsive forces repulsive forces
Electron proton attractive forces attractive forces
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Bond Length Diagram Net repulsionNet attraction Scientists can
determine the internuclear distances that correspond to the lowest
energy states of molecules
http://ch301.cm.utexas.edu/simulations/bond-strength/BondStrength.swf
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Bond Length and Energy Bonds between elements become shorter
and stronger as multiplicity increases
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Covalent Bonds A covalent bond forms when electron pairs are
shared between two atoms whose difference in electronegativity is
1.9 or less an example is the formation of a covalent bond between
two hydrogen atoms the shared pair of electrons completes the
valence shell of each hydrogen.
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Polar Covalent Bonds In a polar covalent bond - the more
electronegative atom has a partial negative charge, indicated by
the symbol - + the less electronegative atom has a partial positive
charge, indicated by the symbol + in an electron density model red
indicates a region of high electron density blue indicates a region
of low electron density
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Polar and Nonpolar Molecules ammonia and formaldehyde are polar
molecules acetylene is a nonpolar molecule
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Carbon Intro and Review Atomic Structure Atoms made up of
protons, neutrons, electrons Isotopes same # protons; different #
neutrons Electronic Structure Electrons determine structure give
rise to bonding behave like waves orbitals (s, p)
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Orbital overlap to form bonds.
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Orbital overlap to form bonds.
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Electron Probabilities and the 1s Orbital The 1s orbital looks
very much like a fuzzy ball, that is, the orbital has spherical
symmetry The electrons are more concentrated near the center
Spherical symmetry; probability of finding the electron is the same
in each direction. The electron cloud doesnt end here the electron
just spends very little time farther out.
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Electron Probabilities and the 2s Orbital The region near the
nucleus is separated from the outer region by a spherical node - a
spherical shell in which the electron probability is zero The 2s
orbital has two regions of high electron probability, both being
spherical
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The Three p Orbitals 2p
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The Five d Orbitals 3d
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Rules for Electron Configurations Subshell filling order...
Each subshell must be filled before moving to the next level 1s 2
2s 2 2p 6 3s 2 3p 6... 1s < 2s < 2p < 3s < 3p < 4s
< 3d < 4p < 5s < 4d < 5p < 6s
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The most stable arrangement of electrons in subshells is the
one with the greatest number of parallel spins (Hunds rule). C 1s 2
2s 2 2p 2 N 1s 2 2s 2 2p 3 O 1s 2 2s 2 2p 4 F 1s 2 2s 2 2p 5 Ne 1s
2 2s 2 2p 6
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Periodic Relationships The valence shell is the outermost
occupied shell The period number = principal quantum number, n, of
the electrons in the valence shell
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Atomic Orbitals 1s 1 st orbital s type (spherical) 1s, 2s,
3s
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Atomic Orbitals 2s orbital (spherical)
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Atomic Orbitals p (2p, 3p) 3 orbitals oriented perpendicular to
each other have node (region of 0 e - density) nodal plane 2p
orbital
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Atomic Orbitals p (2p, 3p) 3 orbitals oriented perpendicular to
each other have node (region of 0 e - density) nodal plane shape
dumbbell
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Chapter 1 Electronic Configuration of Atoms Aufbau Fill lowest
energy orbital 1 st Hunds Rule 1 e - into each orbital of = energy
Pauli Exclusion Principle Pauli Exclusion Principle Electrons in
the same orbital are spin paired Electrons in the same orbital are
spin paired
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Electronic Configurations
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WHY DO HYBRIDS ?? 1. Electron pair repulsions are minimized (=
lower energy) 2. Stronger bonds (= lower energy) are formed 3.
Hybrids have better directionality for forming bonds
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Shapes of Atomic Orbitals All s orbitals have the shape of a
sphere, with its center at the nucleus of the s orbitals, a 1s
orbital is the smallest, a 2s orbital is larger, and a 3s orbital
is larger still
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Shapes of Atomic Orbitals A p orbital consists of two lobes
arranged in a straight line with the center at the nucleus
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Orbital Overlap Model A covalent bond forms when a portion of
an atomic orbital of one atom overlaps a portion of an atomic
orbital of another atom in forming the covalent bond in H-H, for
example, there is overlap of the 1s orbitals of each hydrogen
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Hybrid Orbitals We will study three types of hybrid atomic
orbitals sp 3 sp 3 (one s orbital + three p orbitals give four sp 3
orbitals) sp 2 sp 2 (one s orbital + two p orbitals give three sp 2
orbitals) sp sp (one s orbital + one p orbital give two sp
orbitals) Overlap of hybrid orbitals can form two types of bonds,
depending on the geometry of the overlap bonds bonds are formed by
direct overlap bonds bonds are formed by parallel overlap
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sp 3 Hybrid Orbitals Each sp 3 hybrid orbital has two lobes of
unequal size The four sp 3 hybrid orbitals are directed toward the
corners of a regular tetrahedron at angles of 109.5
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sp 3 Hybrid Orbitals orbital overlap bonding in water, ammonia,
and methane
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sp 2 Hybrid Orbitals An sp 2 hybrid orbital has two lobes of
unequal size the three sp 2 hybrid orbitals are directed toward the
corners of an equilateral triangle at angles of 120 the
unhybridized 2p orbital is perpendicular to the plane of the three
sp 2 hybrid orbitals
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a carbon-carbon double bond consists of one sigma ( ) bond and
one pi ( ) bond sp 2 Hybrid Orbitals
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a carbon-oxygen double bond also consists of one sigma ( ) bond
and one pi ( ) bond
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sp Hybrid Orbitals Each sp hybrid orbital has two lobes of
unequal size the two sp hybrid orbitals lie in a line at an angle
of 180 the two unhybridized 2p orbitals are perpendicular to each
other and to the line through the two sp hybrid orbitals
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sp Hybrid Orbitals a carbon-carbon triple bond consists of one
sigma ( ) bond and two pi ( ) bonds
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Hybrid Orbitals Summary of orbitals and bond types
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Examples of sigma bonds formed from sp 3 hybrid orbitals
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Orbital overlap to form bonds.
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Orbital overlap to form bonds.
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..
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Examples of natural acyclic compounds, their sources (in
parentheses), and selected characteristics
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Examples of natural heterocyclic compounds having a variety of
heteroatoms and ring sizes.
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Examples of natural carbocyclic compounds with rings of various
sizes and shapes.
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Isomerism The Molecular Formula of a substance gives the number
of different atoms present. The Structural Formula indicates how
those atoms are arranged. Isomers are molecules with the same
number and kinds of atoms but different arrangements of the atoms.
Structural (or Constitutional) isomers have the same molecular
formula but different structural formulas.
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Constitutional Isomerism the potential for constitutional
isomerism is enormous World population is about 6,000,000,000
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74 Condensed Structural Formulas
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75 Cyclic Molecules
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76 Bond-line Formulas
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77
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In this representation, bonds that project upward out of the
plane of the paper are indicated by a wedge, those that lie behind
the plane are indicated with a dashed wedge, and those bonds that
lie in the plane of the page are indicated by a line. 78
Three-Dimensional Formulas
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writing structural Formulas In a continuous chain, atoms are
bonded one after another. In a branched chain, some atoms form
branches from the longest continuous chain.
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Abbreviated Structural Formulas
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Formal Charge Here, some molecules one or more atoms maybe
charged +ve or ve which comes from the chemical reactions. Its
important to know how to tell where the charge is located.
H3O+H3O+H3O+H3O+
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Formal Charge minus The formal charge on an atom in a
covalently bonded molecule or ion is the number of valence
electrons in the neutral atom minus the number of covalent bonds to
the atom and the number of unshared electrons on the atom.
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Resonance Resonance structures Resonance structures of a
molecule or ion are two or more structures with identical
arrangements of the atoms but different arrangements of the
electrons. resonance hybrid If resonance structures can be written,
the true structure of the molecule or ion is a resonance hybrid of
the contributing resonance structures.
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Resonance
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Physical measurements tell us that none of the foregoing
structures accurately describes the real carbonate ion. 1.31 1.20
1.41 Experimentally, It was found that all three carbonoxygen bond
lengths are identical: 1.31 . This distance is intermediate between
the normal C=O (1.20 ) and C-O (1.41 ) resonance hybrid The real
carbonate ion has a structure that is a resonance hybrid of the
three contributing resonance structures
